1. Chapter 6:
Chemical Bonding
1. Use the periodic table to infer the
number of valence electrons in an
atom and draw its electron dot
structure.
2. Be able to explain the types of bonds
that atoms can form.
3. List the characteristics of the different
types of chemical bonds.
4. Define the vocabulary words.
5. Use electronegativity values to
classify a bond
1

2. Valence Electrons
 Electrons in the highest occupied energy level
of an element’s atoms
 For representative elements, the number of
valence electrons is the same as the group
number of that element (Page 414)
 Shown in electron dot structures
3 5
2 Symbol 1
of the
6 element 8
4 7
Right, left, top, bottom (1,2,3,4)
Then 12 o’clock and counterclockwise (5,6,7,8)
2

3. Valence Electrons (cont’d)
 Electrons in the highest occupied energy level of an
element’s atoms
 Can be figured out using the group numbers in the periodic
table.
 Ex: The elements of Group 1A (hydrogen, lithium,
sodium, etc.) all have a valence number of 1, which
means there is 1 electron in the highest occupied energy
level. The elements of group 7A (fluorine, chlorine,
bromine, etc.) have 7 electrons in the outer energy level.
 The valence numbers also tell us the likely oxidation
state of that element. More on this later.
3

4. Oxidation States
The oxidation state of an atom is the charge it has when it gains or loses
electrons to form it’s most stable electron configuration.
Valence Number Oxidation State
(charge on the ion)
1 +1
2 +2
3 +3 } CATIONS
5 -3
6 -2
}
ANIONS
7 -1
4

5. The Octet Rule
 Gilbert Lewis used this to explain why atoms form
certain kinds of ions and molecule.
 In forming compounds, atoms tend to achieve the
electron configuration of a noble gas (8 valence e -)
 Recall that each noble gas (except He) has 8 electrons
in its highest energy level and a general electron
configuration of ns2 np6
 Exceptions: Molecules with an odd number of
electrons, more than an octet (PCl5), and less than
an octet (very rare)
Example: NO2 has seventeen valence electrons
[Nitrogen contributes five and each oxygen
5
contributes 6 (2 x 6 =12)]

6. The Octet Rule
 An atom’s loss of an electron produces a cation, or
positively charged ion. The most common cations are those
produced by the loss of valence electrons from the metals,
since most of these atoms have 1-3 valence electrons.
Let’s look at sodium (a 1A metal) as an example:
-e-
Na 1s22s22p63s1 Na+ 1s22s22p6
6

7. Practice Problems
 Write the electron dot structure for each of the
following:
1. Na
2. Al
3. N
4. S
5. Kr
6. Chloride ion
7. Oxide ion
 Refer to pages 414, 417, and 418 for answers.
7

8. Practice Problems
Please write the oxidation numbers of the following:
1) Na 11) Po
2) Al 12) Ga
3) F 13) Cr
4) Cl 14) N
5) Mg
6) P
7) Ca
8) Sb
9) I
10) Sc
8

9. Common Polyatomic Ions
Hydroxide: OH- Permanganate: MnO4-
Bicarbonate: HCO3- Ammonium: NH4+
Carbonate: CO32- Acetate: C2H3O2-
Sulfate: SO42- Hydrogen-
Sulfite: SO32- Phosphate: HPO42-
Phosphate: PO43- Dichromate: Cr2O72-
Perchlorate: ClO4-
Nitrate: NO3-
Nitrite: NO2-
Chlorate: ClO3-
Cyanide: CN- 9

10. Chemical Bonding
 Chemical energy & potential energy stored
in chemical bonds
 Atoms prefer a low energy condition
 Atoms that are bonded have less energy
than free atoms- more stable.
 To combine atoms: energy is absorbed
 To break a bond: energy is released (AB)
10

11. Chemical Bonds
 Created when two nuclei simultaneously
attract electrons
 When electrons are donated or received,
creating an ion (anion, cation)
 In most elements, only valence electrons
enter chemical reactions
 Atoms of everyday substances are held
together by chemical bonds (water, salt
anti-freeze)
11

12. Types of Chemical Bonds
1) Ionic Bond: chemical bonding that results from
the electrical attraction between cations and anions
where atoms completely give their electron(s) away
12

13. Types of Chemical Bonds
2) Covalent Bond: chemical bonding that results
from the sharing of electron pairs between two atoms.
The electrons are “owned” equally by the two atoms.
13

14. Relative Forces of
Attraction
 Ability of a nucleus to hold its valence
electrons (Group 7A has a greater ability
to hold on to its valence electrons than
Group 1A)
 Ionization energy: energy required to
lose an electron (As atomic number
increases down a group, the most
loosely bound electrons are more easily
removed, so ionization energy
decreases. For the most part, it
increases along each period.) 14

15.  Electron affinity – tendency to gain an
electron (Energy is released)
 Electronegativity – measure of the
electron attracting power of an atom
when it bonds with another atom
* Fluorine (4.0) is the highest
* Cesium (0.7) is the lowest – least
ability to attract bonding electrons and
thus the greatest tendency to lose an
electron
* Noble gases are not assigned
electronegativities because these
elements do not generally form bonds
15
(inert)

16.  The periodic trend of the
electronegativities is the same as that of
the ionization energies. Thus, as the
atomic number increases along a period,
the electronegativity increases. As the
atomic number increases down a group,
the electronegativity decreases.
 In general, metals have a low
electronegativity and nonmetals have a
high electronegativity
16

17. Electronegativity and Bond Types
Covalent Bonds: bonding between elements with an electro-
negativity difference of 1.7 or less.
Nonpolar-Covalent Bonds: covalent bond in which electrons are
shared evenly by the bonded atoms with an electronegativity
difference of 0 – 0.3.
Polar-Covalent Bonds: covalent bond in which the bonded atoms
have unequal attraction of the shared electrons, and have an
electronegativity difference of 0.4 – 1.7
Ionic Bonds: bonding due to difference in electric charge of two
elements due to loss/gain of electrons. Must have an electro-
negativity of 1.8 – 4.0.
17

18. Electronegativity and Bond Types
Water is a polar molecule, because the electrons
are not shared evenly by the hydrogen and oxygen.
Ionic Bonds: bonds in which electrons are donated from one
atom to another and have an electronegativity difference
of 1.8 or higher.
18

19. Electronegativity and Bond Types
Using the electronegativity values found on page 161 of
your book, predict the types of bonds the following will form.
1) O2
2) NaCl
3) N2
4) Knowing that the electronegativity of sulfur is 2.5, what type
of bond will sulfur form with:
a) hydrogen
b) cesium
c) chlorine 19

20. Electronegativity
Electronegativity is a measure of how strongly an element
can remove an electron from another element.
20

21. Ionic (Electrovalent)
 Bonds chemical bond
The strongest
Complete transfer of electron(s) from one element to another
 Generally formed when metals combine with
nonmetals (Groups 1-2a w/ 5-7a)
 Coulombic forces – electrostatic force in which two
oppositely charged ions are mutually attracted
 Usually occurs when the difference in
electronegativities is 1.8 or greater
21

22. NaCl – Ionic Bond
 Draw
22

23. Writing Ionic
Compounds
 Beryllium fluoride
 Calcium oxide
 Scandium sulfide
 Aluminum chloride
23

24. Ionic Solids
 Form crystal lattice (orderly, repeating,
three-dimensional pattern)
 The charges and relative sizes of the
ions determines the crystal structure
 The number of ions of opposite charge
that surround the ion in a crystal is called
the coordination number of the ion.
24

25.  Poor conductors of electricity (no free
electrons)
 High melting point
 High boiling point
 Brittle and break easily under stress
 Liquid or aqueous: good conductors of
electricity but ionic bond is dissolved
25

26. The Normal
Arrangement of an
Ionic Crystal
- + - +
+ - + -
- + - +
+ - + -
- + - +
Opposite charges attract
26

27. Arrangement when Stress
is Applied
- +
+ - - +
- + + -
+ - - +
- + + -
- +
Adjacent to ions with same charge (repulsion)
27

28. Crystal Lattice is
Destroyed
- +
+ - - +
- + + -
+ - - +
- + + -
- +
Crystal melts, vaporizes, or dissolves in water
(ions free to move about)
Cleavage – splitting along a definite line
28

29. Covalent Bonding
 Electrons are shared
 One atom does not have enough pull on the
electron to take it completely from the other
atom
 Occurs when electronegativity difference is
less than 1.8
 Covalently Bonded Solids:
1. Softness
2. Poor conductor of electricity and heat
3. Low melting point 29

30. Lewis Structures
 Single covalent bond – one shared pair
of electrons:
 H· + ·H H : H or H H
 Double covalent bond – two shared pairs
of electrons :
:
:
:
:
:
: O. + : O. : O: : O: or : O O:
.
.
 Triple covalent bond – three shared pairs
of electrons
:
:
. N. + . N. : N : : : N : or : N N:
. .
30
Note: all of these obey the octet rule

31.  Coordinate covalent bond – one atom
contributes both bonding electrons
 NH3 + H+ [NH4]+
ammonia hydrogen ion ammonium ion
H +
:
H: N : H + H+ H: N : H
:
:
H H
The structural formula shows an arrow
that points from the atom donating the
electrons to the atom receiving them.
Refer to page 444
31

32. How to Construct Lewis Structures
Step 1: Determine the type and # of atoms in molecule
CH3I
has 1 Carbon, 3 Hydrogens and 1 Iodine
Step 2: Write electron dot notation for each type of atom
.. .
. .. .
.. .
C H· I
.
Step 3: Determine the total # of electrons available in
the atoms to be combined.
C 1 x 4e- = 4e-
I 1 x 7e- = 7e-
H 3 x 1e- = 3e-
32
14 e-

33. How to Construct Lewis Structures
Step 4: Arrange the atoms to form a skeleton structure
for the molecule. Then connect the atoms by
electron-pair bonds.
H
. .. .
. . ..
HC I
H
Step 5: Add unshared pairs of electrons to each non-
metal atom so that each is surrounded by 8.
H
. .. .
. . ..
. .
. ...
HC I
33
H

34. Electron Dot Practice: Compounds
1) H2O 5) CCl2H2
2) H2O2 6) NH3
3) HCN 7) N2
4) AlF3 8) CO2
34

35.  A single water molecule is a good example of covalent bonding
between atoms. The hydrogen atoms “share” their electrons with
the larger oxygen atom so that oxygen now has a full outer level
with 8 electrons and each hydrogen has a full outer level with 2
electrons. Oxygen has a higher electronegativity than hydrogen,
so there is actually an uneven sharing of electrons, resulting in a
polar molecule. More on this later.
e- e-
e- e-
shared electrons 8p+ shared electrons
e
-
e-
8n0
e- e-
e- e-
1p+ 1p+
35

36.  Bond dissociation energy: total energy required to
break the bond between two covalently bonded atoms
(remember that energy is measured in joules or
kilojoules)
H–H + 435 kJ H . + .H
 Resonance Structures: refers to bonding in molecules
or ions that cannot be correctly represented by a
single Lewis structure.
36

37. Bond Length vs. Bond
Energy
There is a correlation between bond length and the amount
of potential energy stored in that bond. For example:
Bond Bond Length (pm) Bond energy (Kj/mol)
C C 154 346
C C 134 612
C C 120 835
C N 147 305
C N 132 615
C N 116 887
N N 145 163
N N 125 418
37
N N 110 945

38.  Molecular orbitals – when two atoms combine and their
atomic orbitals overlap
 Sigma bond - molecular orbital that is symmetrical
along the axis connecting two atomic nuclei
In both of these examples, the p orbitals are overlapping
and sharing electrons.
38

39.  pi bond – weaker than sigma bond; usually
sausage-shaped regions above and below the
bond axis (Page 445)
39

40. Examples of Sigma and Pi bonds
H3C – CH3
H2C = CH2
HC – CH
–
–
40

41. VSEPR Theory (page 200)
VSEPR Theory (Valence Shell Electron-Pair Repulsion theory):
states that repulsion between the sets of valence-level electrons
surrounding an atom causes these sets to be oriented as far
away from each other as possible, thus determining the shape
of molecules.
41

42. VSEPR Theory
So then why is H2O bent, but BeF2 is linear?
H2O BeF2
The answer is the free electron pairs. Oxygen has 2 pairs,
beryllium has none.
1
:
: O. . Be.
2
.
42

43. VSEPR Shapes
Linear
Trigonal-Planer
Bent/Angular
Tetrahedral
Trigonal-Pyramidal
Trigonal-Bipyramidal
Octahedral
(#s 3, 5 and 7 are coordinate covalent bonds!)
43

44. VSEPR Theory
These free electron pairs repel each other because they have a
negative charge, and so they force those atoms that are
.
covalently bonded to be pushed as far away as possible.
.
. .
44

45.  Hybridization – several atomic orbitals
mix to form the same total number of
equivalent hybrid orbitals (CH4 – Page
457 )
*Note: An sp3 orbital is an example of a
hybrid.
45

46. Types of Covalent
Bonding
 1. Nonpolar – when atoms have the same or
similar electronegativity; when the atoms in
the bond pull equally and the bonding
electrons are shared equally.
(Generally a difference of 0.0- 0.4)
* Examples: Diatomic elements
(H2 , N2 , O2 , F2 , Cl2 , I2 , Br2)
* Nonpolar Covalent:
Bonded Hydrogen Atoms
46

47.  2. Polar – unequal sharing of electrons
* Pairing of atoms when one has a stronger
attraction for the electrons
* Most compounds are polar covalent
Examples: H2O , NH3 , HF , HCl
*Polar covalent also called dipoles
*Creates partial charges
Partially +
Partially -
Example: HCl (0.9 difference of the electro-
negativities) H (2.1) Cl (3.0)
Electronegativity difference is less than 1.7
47

48. Example: water.
An uneven distribution of the electrons results because
the oxygen has a higher electron affinity than the
hydrogens. Thus, you have a negative and positive end
of the molecule: polarity. Because this molecule has 2
poles, it is called a dipole molecule.
δ-
δ = delta or e- e-
overall Oxygen
e -
e-
e - 8p+ e-
8n0
e- e-
Hydrogens
e- e-
1p+ 1p+
δ+ 48

49. Attractions Between Molecules
Molecules are often attracted to each other by a variety of
forces. The intermolecular attractions are weaker than either
an ionic or covalent bond. These attractions are responsible
for determining whether a molecular compound is a gas, liquid,
or solid at a given temperature. Here is a list of these various
attractions:
1. van der Waals forces: weakest type of intermolecular
attractions.
Dispersion (London) forces: weakest of all molecular
interactions, caused by the motion of electrons.
Increases as the # of electrons increases.
Ex: Cl & F are gases at STP; Br is liquid at STP;
I is solid at STP.
49

50. 2. Dipole Interactions: attraction of polar molecules to one
another. Remember that polar molecules are like magnets;
they have a positive and negative end.
A glucose molecule
in water has many dipole
interactions since both
water and glucose are
polar. The positive poles
of the water molecule are
attracted to the negative
poles on the glucose and
vice versa.
50

51. 3. Hydrogen Bonds: attractive forces in which a hydrogen
covalently bonded to a very electronegative atom is also
weakly bonded to an unshaired pair of electrons on another
electronegative. Hydrogen bonding always involves
hydrogen. Hence the name. Duh.
The hydrogen bonding between
water molecules dictates many of
the properties of water. It also
explains why water is a liquid
rather than a gas at room
temperature.
51

52. Network Solids
 Macromolecules
 Covalent network of atoms bonded
 Absence of molecules throughout the
solid
 Properties
1. Hardness
2. Poor conductor of electricity (electrical
insulation)
3. Poor conductor of heat 52

53.  Examples: diamond graphite (carbon)
Does not melt - vaporizes to a gas at 3500 °C
Carbon atom
Covalent bond
 Boron nitride (BN), asbestos, silicone carbide
(SiC, grindstones), silicone dioxide (SiO2 ,
quartz)
53

54. Metallic Bonds
 Most metallic elements, except liquid
mercury, are solids at room temperature
and exhibit a crystal structure (zinc)
 Arrangement of stationary positive metal
ions surrounded by a “sea of mobile
electrons””
- - - -
+ + +
- - - -
- + - + - + -
- - - -
+ + +
54

55.  Properties:
1. Malleability – ability to be hammered
into different shapes
2. Ductility – ability to be drawn into wire
3. Conductor of heat
4. Conductor of electricity
5. Luster – shine
6. Tenacity – structural strength
(resistance to being pulled apart)
55

56. Alloys
 Mixtures composed of two or more elements, at least
one of which is a metal
 Properties usually superior to those of the component
elements
 Sterling silver – silver and copper
 Bronze – copper and tin
 Steel – iron, carbon, boron, chromium, manganese,
molybdenum, nickel, tungsten, vanadium (Interstitial
alloy)
 Interstitial alloy – smaller atoms fit into spaces between
larger atoms
 Substitutional alloy – atoms of the components are
about the same size (They can replace each other in
the structure.)
56

57. Summary : Types of
Bonds
1. Ionic – complete transfer of electrons
2. Covalent – share electrons
A. Nonpolar : same or similar electronegativity
B. Polar – unequal sharing
Electronegativity Difference:
C < 2.0 ≤ I (Know exceptions)
*Know table on page 465
3. Network solids – covalent network of atoms
(absence of molecules)
4. Metallic – positive ions around a “sea of
mobile electrons”
57

58. Electronegativity
Differences and Bond
Types
 0.0-0.3 Nonpolar covalent
 0.4-1.0 Moderately polar covalent
 1.0-1.8 Very polar covalent
 1.8 or greater Ionic
58

59. General Trends of the
Representative Elements
 Group 1A - lose one electron
 Group 2A - lose two electrons
 Group 3A - lose three electrons
 Group 4A - share, lose or gain 4 e-
 Group 5A - share, gain three electrons
 Group 6A - share, gain two electrons
 Group 7A - gain one electron
 Group 8 - do not react, noble gases
59

60. Think!
 Why is it possible to bend metals but not
ionic crystals?
 In an ionic compound, ions of like charge
do not have mobile electrons as
insulation. When forced into contact by
physical stress, the ions of like charge
repel, causing the crystal to shatter.
60

61. Attr actions
between
Molecules weak
Name and describe the
attractive forces that hold groups of
molecules together.
61

62. Van der Waals Forces
 Weaker than either an ionic or covalent
bond
 Responsible for determining whether a
molecular compound is a gas, liquid, or
solid at a given temperature
 Two types: dispersion forces and dipole
interactions
 Dispersion – caused by motion of
electrons; dispersion generally increases
as the number of electrons increases
62

63.  Example of dispersion forces: (Refer to
Group 7A) F and Cl are gases at STP; Br
is a liquid at STP, and I is a solid at STP
 Dipole interactions – electrostatic
attractions between oppositely charged
regions (Example: water)
63

64. Hydrogen Bonds
 Strongest of the intermolecular forces
 Important in determining the properties of
water and biological molecules such as
proteins
 Has only about 5% of the strength of an
average covalent bond
64